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The ideal gas law takes into account many assumptions. The most important ones are all collisions are elastic, the molecules themselves don't take up space, and there is no interaction between gas particles except for collisions. These assumptions work best at low pressure and at high temperatures. This means the ideal gas law fails at extremely high pressure and extremely low temperatures. The high pressure forces the particles together, and the low temperature caused the particles to slow down. These both cause the particles to interact without touching, decreasing the effectiveness of the ideal gas law.
The ideal gas law fails when the temperature is extremely low and when the pressure is extremely high. When the ideal gas law fails, you can no longer use the ideal gas law equation, you need to use the Van Der Waals Equation, which is quite a bit more complicated than the ideal gas law equation.
Does the Ideal Gas Law only apply to scenarios in Standard Temperature and Pressure? (STP) Or can approximations still be made using the Ideal Gas Law Equation with other temperatures within a close range?
The Van der Waals equation modifies the Ideal Gas Law slightly to better predict the behavior of real gases, taking intermolecular forces and volume taken up by gas molecules into account. It's really interesting to see how this was done but for the purposes of this course it should be safe to assume that all gases are ideal.
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